Riassunto esame Chimica, prof. Saracco, libro consigliato Chemistry: The Molecular Nature of Matter and Change di Silberberg

The components of matter: elements, compounds, and mixtures

Matter can be classified into three types based on composition.

Elements

Elements are the simplest type of matter with unique physical and chemical properties. They consist of only one kind of atom. The macroscopic properties of an element depend on the submicroscopic properties of its atoms. Some elements occur in molecular form: an independent structure of two or more atoms bounded together.

Compounds

A compound consists of two or more different elements that are bonded together chemically. Their atoms have joined in a chemical reaction. Single elements are present in specific parts by mass ratio, since each unit of the compound consists of a fixed number of atoms of each element. For example, ammonia is 14 parts nitrogen by mass and 3 parts hydrogen by mass because an atom of nitrogen weighs 14 times an atom of hydrogen, and a molecule of ammonia has 3 atoms of hydrogen and 1 of nitrogen (NH₃). The properties of a compound are different than those of the elements that form it. A compound can be broken down into simpler substances, its component elements (this is a chemical change). Elements and compounds are both substances with a fixed composition.

Mixtures

A mixture consists of two or more substances that are physically intermingled. It is not a substance because the components of a mixture can vary their parts by mass. A mixture retains many of the properties of its components. A mixture can be separated into its components by physical changes.

Observations that led to an atomic view of matter

Mass conservation

The law of mass conservation states that the total mass of substances during a chemical reaction does not change. The number of substances may change, and their properties must. Based on all chemical experience and this law, we can state that matter cannot be created or destroyed. To be precise, we know from the work of Albert Einstein that mass before and after a reaction is not exactly the same, some is converted to energy or vice versa, but except for nuclear reactions, the changes are small enough to be ignored.

Definite composition

The law of definite composition states that a particular compound is composed of the same elements in the same parts by mass. The fraction by mass is the part that each element contributes. It is obtained by (element mass/compound mass). The percent by mass is the fraction expressed as a percentage.

Multiple proportions

The law of multiple proportions states that if elements A and B react to form two compounds, the different masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers. This means that in two compounds of the same elements, the mass fraction of one element relative to the other changes in increments based on ratios of small whole numbers. For example, in CO, carbon is 0.429 parts by mass and oxygen has 0.571, so their ratio is 1.33. In CO₂ carbon has 0.273 parts by mass and oxygen has 0.727, so their ratio is 2.66. As we can see, 2.66/1.33=2.

Dalton's atomic theory

Postulates of the atomic theory

• All matter consists of atoms that cannot be created or destroyed.
• Atoms of one element cannot be converted into atoms of another element. They can only recombine to form different substances.
• Atoms of an element are identical in mass and other properties and are different from atoms of any other element.
• Compounds result from the chemical combination of a specific ratio of atoms of different elements.

Observations that lead to the nuclear atom model

Discovery of the electron and its properties

Scientists had known matter and electric charge were related, they also knew that an electric current could decompose certain compounds into their elements. To discover the nature of electric current, scientists tried to pass current through an evacuated glass tube fitted with metal electrodes. A ray could be seen striking the phosphorous coated ending of the tube. These rays were called cathode rays since they were emitted by the cathode (-). The main conclusion was that these rays consisted of negatively charged particles present in all matter. Cathode ray particles were named electrons. There were two experiments that determined the mass and charge of the electron. The mass/charge ratio was discovered by J. J. Thompson. He estimated that the mass was less than 1/1000 of the mass of the hydrogen atom, denying Dalton's atomic theory by demonstrating that atoms contained smaller particles. The charge was discovered by R. Millikan, who observed the movement of charged oil droplets in an electric field. He found that all the charges on the droplets were multiples of a minimum charge: -1.602x10^-19, the charge of the electron.

Discovery of the atomic nucleus

Matter is electrically neutral, so atoms must be as well. Thompson had proposed his plum-pudding model that was put to the test by E. Rutherford. In his experiment, he noticed that alpha particles fired against a gold sheet were sometimes deflected. He concluded that the positive charge in the atom was not distributed in a large amount of space, but concentrated in what he called the nucleus. He also concluded that most of the mass of the electron was concentrated in the nucleus. Neutrons (that account for a good part of the atom's mass) were discovered only 20 years later by J. Chadwick.

The atomic theory today

Structure of the atom

An atom is an electrically neutral, spherical entity composed of a positively charged nucleus surrounded by one or more negatively charged electrons. Electrons move in the available volume, held there by the attraction of the nucleus. An atom's diameter (1x10^-10m) is about 20000 times the diameter of the nucleus (5x10^-15m). The nucleus contains 99.97% of the atom's mass. An atomic nucleus consists of protons and neutrons. The proton has a positive electrical charge, the neutron has no charge, the electron has a charge of the same magnitude as the proton, but with the opposite sign.

Atomic number, mass, and symbol

The atomic number (Z) of an element equals the number of protons in the nucleus of each of its atoms. All atoms of an element have the same number of protons, this number is different from that of any other element. The mass number (A) is the total number of protons and electrons in an atom's nucleus. The atomic symbol is the letters through which an element is known (e.g., C for carbon).

Isotopes

All atoms of an element have the same atomic number, but not the same mass number. Isotopes of an element are atoms that have different numbers of neutrons. The chemical properties of an element are primarily determined by the number of electrons, so all isotopes have nearly identical chemical behavior.

Atomic masses of the elements

The mass of an atom is measured relative to the mass of an atomic standard. The amu (atomic mass unit) is 1/12 of the mass of a carbon-12 atom. The value of 1 amu is 1.66054x10^-24g. The isotopic makeup of an element is determined by mass spectrometry. The mass spectrometer also gives the relative abundance as a percentage of the isotopes in a sample of the element. From such data, we obtain the atomic mass of an element. It is the average of the masses of its naturally occurring isotopes weighted according to their abundances. It is an average value; no individual atom of an element has precisely the mass of its element's atomic mass, but we consider a sample to consist of atoms with this average mass.

Elements and periodic table

Organization of the periodic table

Each element has a box that contains its atomic number, symbol, and mass (a mass in parentheses is the mass of the most stable isotope of that element). The boxes lie, from left to right, in order of increasing atomic number. The boxes are arranged into a grid of periods (horizontal) and groups (vertical). Periods have a number from 1 to 7 and groups have a number from 1 to 8 and a letter (A or B). The eight A-groups contain the main-group elements. The 10 B-groups contain the transition elements. The two horizontal series of the inner transition elements (lanthanides and actinides) fit between group 3B and 4B. They are placed below the main body of the table.

Classification of the elements

The metals lie in the large bottom left portion of the table. Three quarters of the elements are metals and are generally shiny solids at room temperature. They conduct heat and electricity well and are malleable and ductile. Nonmetals lie in the small upper right portion of the table; they are generally gases or dull, brittle solids at room temperature. They conduct heat and electricity poorly. Metalloids (or semimetals), which lie along the staircase, have properties between those of metals and semimetals.

Important notions

In general, elements in a group have similar chemical properties, while elements in a period have different chemical properties. Despite the classification, there is a gradation in properties from left to right and top to bottom. Group 1A consists of the Alkali metals, 2A consists of the alkaline earth metals, both very reactive. Group 7A, halogens, are highly reactive nonmetals. Group 8A, noble gases, are relatively nonreactive nonmetals.

Compounds and introduction to bonding

The overwhelming majority of elements occur in compounds combined with other elements. Elements combine in two general ways, both involving the electrons of the atoms of interacting elements:

• Transferring electrons from one element to the other to form ionic compounds.
• Sharing electrons between atoms of different elements to form covalent compounds.

These processes generate chemical bonds, the forces that hold atoms together in a compound.

The formation of ionic compounds

Ionic compounds are composed of ions, charged particles that form when an atom (or small group of atoms) gains or loses one or more electrons. A binary ionic compound, the simplest of these, generally forms when a metal reacts with a nonmetal: each metal atom loses an electron and becomes a cation, while each nonmetal atom gains one and becomes an anion. Cations and anions attract each other to form the ionic compound. A cation or anion derived from a single atom is called a monatomic ion. The strength of the bond is dictated by Coulomb's law \((F=(q_1 \times q_2)/r^2)\). This means that ions with higher charges attract each other more, and the smaller the ions, the greater the force (less distance).

Number of electrons lost or gained

Ionic compounds are neutral; there is an equal number of Na⁺ and Cl^- ions in sodium chloride. For A-group elements, we usually find that the metal atoms lose electrons, and nonmetal atoms gain them to form ions with the same number of electrons as in an atom of the nearest noble gas on the periodic table (noble gases are very stable thanks to the number and arrangement of their electrons). Metals lose electrons according to their group (1A loses 1, 2A loses 2, etc.). Nonmetals gain electrons according to their group (7A gains 1, 6A gains 2, etc.).

The formation of covalent compounds

Covalent compounds form when atoms share electrons, usually between nonmetals. Two H atoms approach each other: as they get closer, the nucleus of one attracts the other's electron and vice versa. As they start to interpenetrate each other, repulsions increase until an optimum point is reached. This is when a covalent bond forms: a pair of electrons mutually attracted by two nuclei. A molecule of H₂ is formed. In the same way, atoms of different elements share electrons to form covalent compounds.

Distinguishing the entities in covalent and ionic substances

There is a key distinction between the chemical entities in covalent substances and in ionic substances. Most covalent substances consist of molecules, but there are no molecules in an ionic compound. It is a continuous array in three dimensions of oppositely charged ions. Another key distinction concerns the nature of the particles attracting each other. Covalent bonding involves the mutual attractions between two nuclei and two electrons. Ionic bonding involves the mutual attraction between positive and negative ions.

Polyatomic ions: covalent bonds within ions

Many ionic compounds contain polyatomic ions, that consist of two or more atoms bonded covalently that have a net positive or negative charge.

Stoichiometry of formulas and equations

The mole

The mole, abbreviated mol, is the SI unit for the amount of substance. It is defined as the amount of substance that contains the same number of entities as the number of atoms in 12g of Carbon twelve. This number is called Avogadro's number (6.022x10²³). The mole gives you both the number of entities in a given substance and its mass.

• Elements: The mass in amu of one atom of an element is the same numerically as the mass in grams of one mole of atoms of that element.
• Compounds: The mass in amu of one molecule of a compound is the same numerically as the mass in grams of one mole of the compound.

Determining molar mass

The molar mass (M) of a substance is the mass per mole of its entities and has units of grams per mole.

• Elements: Look up the atomic mass and note whether the element is monoatomic or molecular. If it is monoatomic, the molar mass is the periodic table value in grams per mole. If it is molecular, you must know the formula to determine molar mass.
• Compounds: The molar mass is the sum of the molar masses of the atoms in the formula.

Converting between amount, mass, and number of chemical entities

• Converting between amount and mass: \(m=(\text{number of moles})\times(\text{mass per mole})\)
• Converting between amount and number: \((n \text{ of entities})=(\text{number of moles})\times N_a\)
• To solve mass-number conversion involving compounds, one must know the chemical formula.

The importance of mass percent

It is important to know how much of an element is present in a given amount of a compound. We find the composition in terms of mass percent and use it to find the mass of each element in the compound.

• For a molecule of a compound: \((\text{mass % of x})=((\text{atoms of x in the formula})\times(\text{atomic mass of x}))/ (\text{molecular mass of the compound})\)
• For a mole of compound: \((\text{mass % of x})=((\text{moles of x in formula})\times(\text{molar mass of x}))/(\text{mass of 1 mol of compound})\)

Gases and the kinetic-molecular theory

The gas laws and their experiments and foundations

The behavior of a sample of gas can be completely described by four variables: P, V, n, T. They are interdependent, which means that any one of them can be determined by knowing the other three. Because gas volume is so easy to measure, the laws are expressed as the effect on gas volume of a change in gas pressure, temperature, or amount. The ideal gas law describes the behavior of an ideal gas, one that exhibits linear relationships among P, V, n, T. No ideal gas actually exists, but most simple gases and noble gases behave nearly identically even at ambient temperature.

Boyle's Law

The product P*V at constant T and n is a constant. V is directly proportional to 1/P. At a constant temperature, the volume occupied by a fixed amount of gas is inversely proportional to the external pressure.

Charles' Law

V/T is a constant at constant P and n. V is directly proportional to T. At constant pressure, the volume occupied by a fixed amount of gas is directly proportional to its absolute temperature (K°).

Other relationships based on Boyle's and Charles' laws

P/T is a constant at constant T and n. P is directly proportional to T. At constant volume, the pressure exerted by a fixed amount of gas is directly proportional to the absolute temperature.

Avogadro's Law

V/n is a constant at constant T and P. V is directly proportional to n. At fixed temperature and pressure, the volume occupied by a gas is directly proportional to the amount (mol) of gas. This is another way of expressing that at a fixed temperature and pressure, equal volumes of any ideal gas contain the same number of particles (mol).

Gas behavior at standard conditions

Chemists have assigned a baseline set of standard conditions called standard temperature and pressure (STP). These are:

• 0°C (273K°)
• 1 atm (760 torr, 100 kPa)

Under these conditions, the volume of 1 mol of an ideal gas is called the standard molar volume. It is 22.4 L (0.0224 m³). At STP, He, N₂, O₂, and other simple gases behave nearly ideally.

The ideal gas law

The ideal gas law is PV=nRT. Where R is the universal gas constant (8.314 J/(mol*K)). R has a different numerical value when different units of measure are used (use SI units).

The partial pressure of each gas in a mixture of gases

The gas behavior we have discussed so far was observed in experiments with air, which is a mixture of gases: the ideal gas law holds for virtually any gas at ordinary conditions, whether pure or a mixture.

• Gases mix homogeneously in any proportions.
• Each gas mixture behaves as if it were the only gas present (assuming no chemical interactions).

Dalton's law of partial pressures

Each gas in a mixture exerts a partial pressure equal to the pressure it would exert by itself. Dalton's discovery was that in a mixture of unreacting gases, the total pressure is the sum of the partial pressures of the individual gases \((P_{\text{tot}}= P_1+P_2+P_3+...+P_n)\). Each component in a mixture contributes a fraction of the total number of moles in the mixture, this portion is the mole fraction (X) of that component. Multiplying X by 100 gives the mole percent.

The ideal gas law and reaction stoichiometry

Many reactions involve gases as reactants or products. From the balanced equation for such a reaction, you can calculate the amounts (mol) of reactants and products. You use the ideal gas law to convert between gas variables and amounts (mol) of gaseous reactants and products.

The kinetic-molecular theory: a model for gas behavior

The kinetic-molecular theory is the model that accounts for macroscopic gas behavior at the level of individual particles. Developed by J. K. Maxwell and L. Boltzmann, the theory draws conclusions based on a few postulates.

Questions concerning gas behavior

Observing gas behavior at the macroscopic level, we must derive a molecular model that explains it.

• Origin of pressure: Pressure is a measure of the force exerted by gas particles colliding with the walls of their container.